Module 2.2 Flashcards

(78 cards)

1
Q

Define an orbital:

A

A region of space 2 electrons can occupy, region with the highest probability of electrons being there.

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2
Q

What shape are S orbitals?

A

Spherical

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3
Q

What shape are P orbitals?

A

8 shape

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4
Q

Define isoelectric:

A

Same electron config as a different element

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5
Q

Define ‘ground state’:

A

All the electrons are in the lowest energy level.

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6
Q

Define ‘excited’:

A

Electrons in a higher that expected energy lever, further out shell.

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7
Q

Describe afbaus theory:

A

Electrons occupy lowest possible energy levels, want to be in ground state.

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8
Q

Describe Hunds rule of multiplicity:

A

Electrons will only occupy an occupied orbital if there are no others, still following law of filling.
Electrons sharing orbitals will have an opposite spin.

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9
Q

Which to elements don’t follow the pattern of the periodic table electron configuration?

A

Chromium and copper

Cr 4s1 3d5
Cu 4s1 3d10

Because they are more stable.

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10
Q

Why don’t transition metals follow the periodic table electron config pattern?

A

When positive ions are formed, 4s electrons are lost first.

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11
Q

Define covalent bond:

A

Electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms.
Non metals bonded to non metals.

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12
Q

Define ionic bonding:

A

Electrostatic attraction between oppositely charged ions.

Metals bonded to non metals

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13
Q

Name for a metal ion?

A

Cation

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14
Q

Name for a non metal ion

A

Anion

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15
Q

Define a metallic bond:

A

Electrostatic attraction between positively charged metal ions and delocalised electrons.

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16
Q

What happens as the charge of an ion increases in a metallic bond?

A

More electrons lost, stronger the metallic bond.

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17
Q

How many electrons can an orbital hold?

A

2

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18
Q

What kind of electron configuration do elements want to achieve by bonding?

A

That of the noble gases

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19
Q

What are dative covalent bonds sometimes called?

A

Coordinate bonds

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20
Q

Describe a dative covalent bond:

A

Covalent bond where both electrons come from same atom.

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21
Q

Describe simple covalent molecules:

A

Small molecules
Don’t conduct electricity
No ions or free electrons
Soluble in organic compounds, not water

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22
Q

Describe the structure of diamond:

A

Carbon covalent ly bonded to 4 others

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23
Q

Describe the properties of diamond:

A

High melting point
Doesn’t conduct electricity, no ions or delocalised electrons
V hard
Insoluble

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24
Q

Describe the structure of graphite:

A

Carbon covalently bonded to 3 others

Delocalised electrons between layers

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25
Ion of aluminium
Al3+
26
Ion of ammonium
NH4+
27
Ion of barium
Ba2+
28
Calcium ion
Ca2+
29
Copper 2 ion
Cu2+
30
Lead ion
Pb2+
31
Silver ion
Ag+
32
Zinc ion
Zn2+
33
Carbonate ion
CO3 2-
34
Hydrogen carbonate ion
HCO3 -
35
Hydroxide ion
OH-
36
Nitrate ion
NO3 -
37
Sulfate ion
SO4 2-
38
Define electonegativity:
The ability for an element to attract the electron pair in a covalent bond to itself
39
Describe a non-polar bond:
Same electronegativity Electrons shared 100% covalent
40
Describe a polar bond:
One pulls electrons closer to its end one is slightly more positive, one is more negative
41
Describe the Pauling scale
Measures electonegativity N O F Cl Br >0.4 difference gives polarity
42
Define a dipole
An uneven distribution of electrons
43
What is the strongest type of intermolecular bond?
Hydrogen bonding | Mainly in water
44
What does a hydrogen bond need?
H atom attached to an electronegative oxygen fluorine of nitrogen O F N must have a lone pair of electrons
45
Boiling point of water and methane
100* C | -164* C
46
Where is the H bond in water?
From the lone pair of oxygen electrons to the electron deficient hydrogen ion
47
3 types of dipole:
Permanent Induced Instantaneous
48
Describe a permenant dipole:
Uneven distribution of electrons | One element in molecule is more electronegative so pulls the electrons to one side
49
Describe an induced dipole:
E.g. Cl-Cl H-Cl The HCl causes the chlorine to have a dipole. Electrons pulled away from Cl on left to Cl on right np y electron deficient hydrogen
50
Describe an instantaneous dipole:
Between elements with same electronegativity Electrons not equally dispersed around atom, maybe one side Electrons randomly move towards one atom No cause May cause an induced dipole in neighbouring molecules More shells=more electrons=larger risk of instantaneous
51
Valence definition
Outer shell
52
2bp
180* linear
53
3pb
120* trigonal planar
54
4bp
109.5 tetrahedral
55
5bp
120* 90* Trigonal by pyramid
56
6bp
90* | Octahedral
57
3bp 1lp
107* | Pyramidal
58
How much do lone pairs reduce bond angles by?
2.5
59
2bp 2lp
Non linear | 104.5
60
Example of a linear molecule
Carbon dioxide
61
Example of a trigonal planar molecule
Boron triflouride
62
Example of a tetrahedral molecule
Methane
63
Boron trifluoride formula
BF3
64
Methane formula
CH4
65
Trigonal by pyrimadol molecule example
Phosphorus pentachloride
66
Phosphorus pentachloride formula
Pcl5
67
Octahedral molecule example
Sulfur hexaflouride
68
Soulful hexaflouride formula
SF6
69
Pyramidal molecule example
Ammonia
70
Ammonia formula
NH3
71
Non linear molecule example
Water
72
3 types of van der waals forces
Permanent dipole - permanent dipole attraction Permanent dipole - induced dipole attraction Instantaneous dipole - induced dipole attraction
73
What are the 4 types of crystalline structure?
Giant ionic structure Simple molecular structure Macromolecular structure Metallic structure
74
Acid + Base =
Salt + water
75
Metal oxide + Acid =
Salt + water
76
Metal hydroxide + Acid =
Salt + Water
77
Metal + Acid =
Metal Salt + Hydrogen
78
Metal carbonate + acid =
Metal salt + carbon dioxide + water